Standardization of NaOH Solution by Colorimetric Titration
Background
Sodium hydroxide is very difficult to accurately measure using a balance because it attracts
water molecules from the surrounding air. Titration is a method of reacting a solution of
unknown concentration (the NaOH in this lab) with one of known concentration. This lab will
use a crystalized acid called potassium hydrogen phthalate (KHP) as the compound of known
concentration. By comparing the amount of NaOH solution needed to completely react a known
amount of KHP, the concentration of the NaOH solution can be determined.
KHP and NaOH react according to the balanced chemical equation below:
1: What is representing NaOH in the reaction above?
2: How many moles of NaOH will react with 0.5 moles
of KHP?
3: What type of reaction is this?
4: KHP’s chemical formula is C8H5KO4. What is the
molar mass of KHP?
Materials
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Potassium hydrogen phthalate (KHP)
~0.1 M sodium hydroxide (NaOH)
Phenolphthalein indicator solution
250 mL Erlenmeyer flask
50 mL buret
Analytical balance
Figure 1
In this titration, you will be placing a known quantity of KHP (dissolved in water) into a flask as
shown in figure 1. You will be adding NaOH to the KHP solution until all of the KHP has been
used up in a neutralization reaction.
5: Is the pH of the KHP solution in the flask acidic or basic?
6: How will we know when all of the KHP is used up?
In this lab, you will conduct a colorimetric titration which means that a color changing
indicator (phenolphthalein) will be used to show changes in the pH. As KHP is reacted with
NaOH, the pH will slowly rise. There will be a point where all of the KHP has reacted and you
will not have added any more NaOH than was needed to react with the KHP. This is called the
equivalence point. At this point, the pH of the solution should be the same as the pH of water
with a salt in it. This is a neutral solution that would have a pH of 7. Phenolphthalein is
colorless in acidic solution and pink in basic solutions. At the equivalence point the solution
should be a very, very faint pink color but the color should be throughout the solution and it
should stay for at least 20 seconds. Eventually the pink color will fade as CO2 is dissolved in the
water making the water slightly more acidic.
7: What should you see in your solution when you have reached the equivalence point?
It is often good to do some pre-planning that will help you later in the lab. If you can start out
with an idea of how much NaOH solution would be needed to react with a given mass of KHP
you should be able to have a better idea of how much of each compound you will be working
with.
8: How much KHP would react completely with 25mL of 0.1M NaOH?
PLEASE CHECK YOUR RESULTS WITH YOUR TEACHER BEFORE MOVING ON
Procedure
Make a table to contain all of the data collected during the lab. (A suggestion ↓)
Trial #
Mass KHP
Initial Volume
NaOH
Final Volume
NaOH
Volume NaOH
Used
1
2
3
4
- From the reagent bench you should get one Erlenmeyer flask and clean it as thoroughly as you
can. [DO NOT USE SOAP as it will leave a residue that is slightly basic that will interfere with
your results.] It is not important to dry the inside of the flask…just get as much water out as
possible.
-In the flask, mass out the amount of KHP calculated in your answer to Question 8 above.
RECORD YOUR ACTUAL MASS IN YOUR DATA TABLE TO 0.001g!!!!! Skipping this will
void the trial.
-Add 30-50 mL of water to the flask and swirl until all of the KHP is dissolved. Check the sides
of the flask to make sure that all KHP crystals are making it into the solution.
-Add 2-3 drops (that’s all you need) of the phenolphthalein solution to the flask and swirl to mix.
-Record the initial volume of NaOH in the buret on your data table. If you skip this step, you
will have to start over!
-Add NaOH from the buret until you reach the equivalence point.
 At the beginning you can add NaOH very quickly. Remember that you calculated
you KHP as enough to react with 25mL of NaOH. Those values were approximate so
they won’t be perfect but they should be close. Gently swirl the flask to mix.
Reminder: The volume of NaOH used is the same as the change in the
volume of NaOH in the flask. Change is always calculated as final-initial.
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When you get close to the equivalence point you want to add NaOH much more
slowly and in much smaller quantities. You can add drop by drop until you are very
close. Gently swirl the flask to mix.
When you feel you are just a few drops from the equivalence point, you should begin
adding fractions of a drop which can be done by just barely opening the stopcock
until a drop is visible and hanging on the buret tip. You can wash the drop into the
flask with a wash bottle filled with DI water. Gently swirl the flask to mix.
9: Why is it ok to add water to the flask if that will change the concentration of the KHP
solution?
When you have reached the equivalence point, record the volume of NaOH in the buret on your
data table. You can use the initial and final volumes to calculate the volume of NaOH used.
10: With only units, write out how you can calculate moles of NaOH reacted from the mass
of KHP.
Repeat these procedures for at least four trials and calculate the molarity of NaOH from each
trial.
In the end, you will need three trials that give calculated NaOH molarities that are within 1.0%
of each other. You will likely need to conduct more trials in order to get three runs that are close
enough together to satisfy this requirement.
Calculations
You will need to calculate the number of moles of NaOH reacted in each of the flasks using the
calculation that you worked out in question 10.
Please calculate the molarity of NaOH for all of your runs. SHOW YOUR CALCULATIONS!
Choose your best three runs and label them.
Put your results from the best three runs in a table like the one below. Calculate the average
molarity of NaOH from the three runs and finish error calculations:
SHOW ALL WORK FOR YOUR CALCULATIONS!!! DO NOT JUST PUT DATA INTO THE TABLE
Trial
NaOH reacted
Molarity NaOH
#
(mol)
(mol/L)
1
2
3
M = moles solute/L solution
Average
Difference
Molarity (mol/L) from Avg.
Percent Error
Difference from Average = |Mrun - Maverage|
Percent Error = (Difference from Avg. / Average Molarity)·100
What to Hand In
Along with answering all questions from the lab instructions, you will need to clearly indicate (in your
tables) all lab measurement data: the mass of KHP, initial and final buret readings etc.
Be sure that you have shown detailed calculations (done by hand!—show your work) for each of your
titrations. The trials that you are “using,” which contribute to the error calculations, should be clearly
identified, but the calculations and data for all of your trials, both good and bad, should appear
somewhere in your lab paper. Be sure to indicate the units of every value.